Unit 3: The Atom

I. History and Development of the Atom

• Believed atoms were:

o The smallest particle of matter

o ___________________________ - could not be

divided or broken down any further

A. Democritus (around 400 B.C.)

indivisible and indestructible

• Based on his observations of the natural world around

him, Democritus was the first to suggest that _________

was _______________________—called them _______all matter

made up of small particles atoms

• Studied the ratios in which elements

combined in chemical reactions.

• Based on his experiments, he formulated

the first real theory about atoms:

Dalton’s Atomic Theory:

1. All _____ is ________ of indivisible particles called _____

B. John Dalton (1803-1805)

2. ___________________________________

(same mass and properties). Atoms of

different elements have different

masses/properties

All atoms of a given element are identical

3. Atoms of elements _______________________ to

_______________; compounds are formed when 2 or more

different atoms bond together

atomsmatter made up

combine in definite ratios

form compounds

4. Atoms _______________________________________

________ - they are just rearranged

• Based on his theories, Dalton viewed the atom as a _______

____________

Dalton’s Atomic Model: ______________

*Note – Dalton didn’t have one specific experiment regarding the atom

Billiard Ball Model

cannot be created or destroyed in a chemical

reaction

hard,

solid sphere

• Performed experiments using a _______________

o Involved shooting a cathode ray (a stream of electricity)

through a tube that had a magnetic field

• Observed two main things:

1. The rays were actually streams of unknown particles that

were so light, they were lighter than the mass of the smallest

known atom (hydrogen)

2. The rays were attracted to the positive plate

• Concluded two main things:

1. The atom really is divisible and it is made up of even smaller

particles

2. One of the particles is negatively charged

C. JJ Thomson (1906)

VIDEO CLIP

cathode ray tube

• He called these __________________________________

(particles “beneath” the atom) ____________

• Based on his experiments, Thomson pictured the atom as a

sphere of positively charged matter with electrons

mixed/embedded in it

Thomson Atomic Model: _________________

VIDEO CLIP

Plum Pudding Model

negatively charged subatomic particles

electrons (e-)

*Plum pudding is a British

dessert. If it helps, think of a

chocolate chip cookie

instead—the chocolate

chips are the electrons and

the dough is the positively

charged sphere

• Performed an experiment called the __________________

o He bombarded (fired) alpha particles (42He)—which are

positively charged—at a thin piece of gold foil

D. Ernest Rutherford ( 1911 )

o If the Thomson model was correct, all the alpha

particles would pass through the foil undisturbed due to

the charge of the positive sphere cancelling out the

negative, free-floating electrons. However, some

particles were slightly deflected

Gold Foil Experiment

• Based on the observation that some alpha particles

were deflected, he concluded that:

1. ______ are made up of ________________

2. There was a ______________________________

_______ = the discovery of the __________!

Rutherford Atomic Model: _____________Nuclear Model

• Based on the observation that most alpha particles

passed through un-deflected, he concluded that:

*Note - Provided no information about ________ other than the fact that

they were located _________________.

VIDEO CLIPhttp://www.mhhe.com/physsci/chemi

stry/essentialchemistry/flash/ruther1

4.swf

Atoms mostly empty space

small, dense, positively charged

center nucleus

electronsoutside the nucleus

E. Neils Bohr (1913)

• Expanded the atomic model by analyzing _________

_____________________

o Emission spectra = a chart of lines of light given off

when an electric current is run through an atom

• Concluded that the __________________was

_______ by the ___________________

o ______________on the spectrum made him

conclude that __________ must be moving from

____________________

the emission

spectra of hydrogen

light emitted/given off

caused movement of electrons

Different colors

electrons

different energy levels

• Based on his experiments, Bohr’s model of the atom had

__________________________________ in well-defined

_________ called _________

o ____________ in ________________had ___________

__________________

• It looked like a solar system – the nucleus was like the sun and

the electrons orbited around the nucleus like planets

Bohr Atomic Model: _____________Planetary Model

electrons traveling around the nucleus

paths orbits

Electrons different orbits different

amounts of energy

F. Werner Heisenberg (1926)

• Bohr’s model only explained the hydrogen atom

with one electron. Did not explain multi-electron

atoms

• Based on his research with multi-electron atoms,

Heisenberg proposed that electrons do not travel

around in circles around the nucleus, instead, they

____________________________

o Orbital: A region in which an electron is most

likely (high probability) located

randomly move in regions (orbitals)

Heisenberg Atomic Model: ___________________

• The nucleus was still the dense, positive center but now it was believed that one cannot know the exact position of an electron; there were only areas where the electron is most likely found (orbitals)

• Heisenberg viewed the atom more like a bee and a hive whereas the Bohr model is like orbiting planets around the sun

Wave-Mechanical Model

* Note - also called the Modern Atomic Model, Quantum-Mechanical Model, or Electron Cloud Model

Practice

1. Which of the following did Rutherford’s Gold Foil experiment prove?

a) That the atom was a uniformly dense sphere.

b) That the atom is mostly empty space with a dense, positive core.

c) That most the atom consists of a uniform positive “pudding” with

small negative particles called electrons embedded throughout.

d) That electrons travel around the nucleus in well-defined paths

called orbits.

2. J.J. Thomson’s Cathode Ray Tube experiment led to the discovery of

a) the positively charged subatomic particle called the electron

b) the positively charged subatomic particle called the proton

c) the negatively charged subatomic particle called the proton

d) the negatively charged subatomic particle called the electron

3. According to the Bohr Model, a) electrons are found in areas of high probability called orbitals

b) electrons travel around the nucleus in circular paths called orbits

c) electrons are found in areas of high probability called orbits

d) electrons travel around the nucleus in random paths called orbitals

4. According to the Wave-Mechanical Model, a) electrons are found in areas of high probability called orbitals

b) electrons travel around the nucleus in circular paths called orbits

c) electrons are found in areas of high probability called orbits

d) electrons travel around the nucleus in random paths called orbitals

II. Atomic Structure

p + 1 1 amu nucleus

n 0 1 amu nucleus

Particles in

nucleus are

called

________

e- - 1 0 amu Outside

nucleus

* Amu = _____________

A.Subatomic Particles

• Subatomic Particles =_____________________

• There are ___of them

particles inside the atom

3

Atomic mass unit

nucleons

Practice1. Which subatomic particle is neutral?

Neutron

2. Where is most of the mass of an atom

located?

the Nucleus

3. What is the charge of the nucleus of any

atom?

Positive

• Atomic Number = __________________________

o Found _______________(the bolded number)

o ________ the number of ________ in an atom

Example: Iron : ________ = ___= _________Atomic # 26

• Nuclear Charge = the ________ of the ________

o The particles in the nucleus are protons and

neutrons. Protons have a charge of +1 and neutrons

have a charge of 0

o Therefore, the nuclear charge is _____________and

_________________________

B. Vocabulary and Notation

26 protons

Example: Carbon – Atomic # =____= ___ protons =

nuclear charge of ____

6 6+6

Vocabulary

identifies the type of element it is

Equals protons

charge nucleus

always positive equal to the number of protons

on periodic table

Example: Cobalt (Co) : _____p and _____ e-27

• Mass Number = The mass of a specific

isotope(sample) of an element

o Mass # = _________________

Why does it make sense that electrons

aren’t included?

o Always a whole number

27

• Atomic Charge = The total charge of an atom

o An atom is ALWAYS _____________

___________ = __________

Example: If an isotope of nitrogen has 7 protons

and 7 neutrons, its mass number is ___

neutral (zero)

# protons # electrons

# protons + # neutrons

So light they barely contribute to mass of element

14

*C – 14 - ____ p: ____ n; ____ e-6 8 6

Mass #

Examples: 94 Be ______ p; ____ n; _____ e-4 5 4

Mass #

Atomic #

• Isotopic Notation: Shows the mass number of an atom

along with element symbol

*C-14 14C Carbon – 14

They all mean the element carbon with a mass number of 14

Notation

Practice

Element Atomic # Mass # Number

of

Protons

Number

of

Neutrons

Number

of

Electrons

Nuclear

Charge

Na - 23

35Cl

K-40

Silver 108

Use your Periodic Table and your knowledge of the atom

to fill in the following chart

11 23 11 1112 + 11

17 35 17 18 17 +17

19 40 19 21 19 +19

47 47 61 47 +47

C. Electrons

• Electrons and how they behave are responsible for many parts of

chemistry

• Even though it is not technically correct, Bohr’s model of the atom is

often used when discussing electrons and the structure of the atom.

It is easiest to visualize and it is “good enough”

• According to Bohr’s model, electrons are located outside of the

nucleus in energy levels. Each energy level can hold a certain

amount of electrons.

Energy level # of electrons

n=1 2

n=2 8

n=3 18

n=4 32

The __________________is from the nucleus, the ___________it

has; therefore, it is ___________and _____________

Closest to nucleus

Furthest away

Energy Levels

farther the electron more energy

less stable easier to move

Electron Configurations

• Electron Configurations = a dashed chain of numbers that shows _______________________________________

o found in the lower left corner of an element box (see below)

• Tells us the number of energy levels as well as the number of electrons in each level

Example: Carbon’s electron configuration is ____

This means it has __ electrons in the ________energy shell and __ in the ________ energy shell (so a total of __ electrons in the atom)

*All electron configurations on the Periodic Table are for atoms when they are most stable (notice #p= #e-)

how electrons are arranged around nucleus

2-4

2 first4 second 6

SUBSTANCE ELECTRON CONFIGURATION

Magnesium

Bromine

*Lead

(see the * at

bottom of

periodic table)

*shortcut allows you to cut out the first two energy levels to shorten

the configuration so it can fit in the box

Practice

Use your Periodic Table to fill in the electron configurations

for the atoms of the following elements

2-8-2

2-8-18-7

On PT = -18-32-18-4

Actually = 2-8-18-32-18-4

Types of electrons

• There are two types of electrons: valence electrons and kernel

electrons

Valence Electrons:

• electrons found in the ___________ shell or energy level

• the ___________in the electron configuration

• the electrons that get lost or gained because they are the furthest away

from the nucleus so they are the easiest to remove

• An element is ___________when its ___________________________

(valence shell)

o __________

o *Hydrogen and Helium are exceptions-stable with 2* Why?

Kernel Electrons:

• Inner electrons (all the other, non-valence electrons)

Example: Calcium’s configuration is 2-8-8-2; therefore it has __

valence electrons and ____ kernel electrons.

outermost

last number

most stable last occupied energy level is full

8 is great!

The first shell IS full with only 2 e-

218

Practice

Electron

configuration

# valence e- # kernel e-

Chlorine

Nitrogen

Sodium

Use your PT to fill in the electron configurations for the atoms of the

following elements. Then identify the # of valence and kernel e-

2-8-7 7 10

2-5 5 2

2-8-1 1 10

D. Atom Diagrams

• There are two common diagrams used to represent the

structure of the atom: Bohr Diagrams and Lewis Dot

Diagrams

Bohr Diagrams

• As previously mentioned, Bohr’s model is often used when

visualizing an atom

• Bohr Diagrams are models of the atom that have the

electrons in rings (orbits) around the nucleus

Steps for drawing Bohr Diagrams:

1. Draw a circle representing the nucleus

2. Find the number of protons and neutrons and write them inside the

nucleus

• To find # of protons-find the element’s atomic number using the

Periodic table

• To find # of neutrons – subtract the atomic number (or number

of protons) from the mass number

3. Look up the element’s electron configuration on the Periodic Table

4. Use the electron configuration to determine how many rings will be

around the nucleus (# of energy levels = # of rings)

Example: Magnesium’s configuration is 2-8-2 so there will be

______________around the nucleus

5. Using dots to represent electrons, fill in the number of electrons in

each ring

3 rings/circles

Example: Draw the Bohr Diagram for C-14

6 p

8 n...

.

.

.

Electron configuration: 2-4

Lewis Dot Diagrams

• Whereas Bohr Diagrams illustrate all the electrons of an

element, lewis dot diagrams or ___________________,

only illustrate the valence electrons

o Valence electrons are often seen as the most important

ones because they are the electrons that are gained or

lost when elements bond to form compounds

• Electron dot diagrams consist of the _______________

surrounded by dots that represent its _______________

electron dot diagrams

element symbol

valence electrons

Examples:

Ca N F

2-8-8-2

Valence e-

Ca

2-5

N

2-7

F

Notice – Put one electron on each side then double up!

Steps to Drawing Lewis Dot Diagrams:1. Write the element’s symbol

2. Find the electron configuration from Periodic Table. The last number

in the configuration is the number of valence electrons

3. Using dots to represent the electrons, place the electrons around

the element symbol, one at a time, starting first at the 12 spot on a

clock. Then add any remaining valence electrons one at a time to

the 3, 6, and 9 spots and then double up if there are more valence

electrons.

• Note: you must add only one electron at a time

because of bonding

o Bonding site = Where there is only a

single electron or an ________________

(lone electrons are open to attach to other

e- and/or easily lost)

2-4

C

Bonding Sites

unpaired electron

Practice1.What is the maximum number of electrons an atom or an ion can have in its

valence shell?

a. 2

b. 4

c. 6

d. 8

*this means that the most dots you can have in a Lewis dot diagram is 8!

2.The number of bonds an atom of an element can form is the same as the

number of

a. electrons in its valence shell.

b. paired electrons in its valence shell.

c. unpaired electrons in its valence shell.

3. Looking back at your Lewis Dot Diagrams, which element can form the

most bonds?

a. Calcium b. Nitrogen c. Fluorine

III. Ions

• Ion = ____________________________(# of protons

DOES NOT EQUAL # of electrons)

• Ions ____________whereas atoms do not!

Example: 2311 Na +1

Atomic # =

Mass # =

Ion Charge =

23

11

+1

# of p =

# of n =

# of e- =

1123 – 11 = 12

(# protons - Ion Charge)

11- (+1) = 10

A.What is an ion?

atom that lost or gained electrons

have a charge

1. Anion =

Negatively charged ion (atom GAINED e-)

2. Cation =

Positively charged ion (atom LOST e-)

Remember:

a CATion is

PAWsitive

• There are two types of ions

Examples:

7 Li + 1 ______ p ________ n ___________e ________

31 P – 3 ______ p ________ n ________ e ___________

79 Se – 2 ______ p ________ n ________ e ___________

19 F – 1 ______ p ________ n ________ e ___________

3

Atomic #Mass # -

Atomic #

7 – 3 = 4

Atomic #-

Charge

3 – (+1) = 2 Lost e-

(cation)

15 16 18 Gained e-

(anion)

34 45 36 gained e-

(anion)

9 10 10 gained e-

(anion)**Think of weight loss – losing weight/electrons is a positive

thing, gaining weight/electrons is a negative thing. It’s

opposite!(when you gain something, it’s negative)

Cation/anion?

Practice

1. When a neutral atom gains an electron, it becomes a

a) negative cation

b) positive cation

c) negative anion

d) positive anion

2. When a neutral atom loses an electron, it becomes

a) negative cation

b) positive cation

c) negative anion

d) positive anion

3. What is the charge on a magnesium ion that has lost two

electrons? _______

4. What is the charge on a fluoride ion that has gained one

electron? _______

5. The chemical symbol Fe+3 represents

a) cation formed as a result of a iron atom losing 3 electrons

b) cation formed as a result of a iron atom gaining 3

electrons

c) anion formed as a result of a iron ion losing 3 electrons

d) anion formed as a result of a iron ion gaining 3 electrons

6. Give the correct chemical symbol for the ion formed when

oxygen gains 2 electrons: ______

+2

-1

O-2

B. Ion Diagrams

• Bohr Diagrams and Lewis Dot Diagrams can also be

used to represent ions

• The steps are the same as the atom except you must

add or subtract electrons from the last number in the

electron configuration

o The last number represents the electrons in the

shell/energy level furthest from the nucleus so they

are the least stable and the easiest to access.

Remember: if it is a __________, you _________

electrons. If it is a ____________, you ____electrons

(opposite!)

positive ion subtract

negative ion add

Examples:

Draw the Bohr Diagram for 40Ca and 40Ca+2

Draw the Bohr Diagram for 19F and 19F-1

• For a lewis dot diagram you must also change the valence

electrons (add or subtract electrons) in the configuration

before doing the diagram. Also, your final diagram must

include _____________________________

Examples:Ex 1: S vs S-2

ADD 2 e- to the 6 that S normally

has in its valence shell.

Ex 2: K vs K+1

REMOVE 1 e- from the valence

shell of K.

*negative ions always end up

with 8 valence e- (8 dots)

*positive ions always end up

with 0 valence e- (0 dots)

brackets and the charge of the ion.

IV. Electron Transitions

A. Ground State vs. Excited State

What do you notice in the diagrams?

• Ground State = Electrons in lowest energy configuration/energy levels possible (____________________________________)

o Stable

• Excited State = Electrons are found in a higher energy configuration (_______________________________)

o Unstable

o excited state electron configuration for Li could be 1-2, 1-1-1 vs. 2-1 ground state

the configuration found on periodic table

any configuration not found on PT

Distinguish between ground state and excited state electron

configurations below:

2-5

2-8-8-1

2-7-1

1-6

Examples:

Ground

Ground

Excited

Excited

• Hint: When atoms are in the excited state, they are still atoms, meaning

protons=electrons. Instead of searching aimlessly for the configuration on

the table, do the following:

1. add up the total number of electrons in the configuration

2. Because it’s an atom, p=e so now that you have the e- you can find the

protons/atomic #/what element it is.

3. Compare the configuration you are given to the one on the table. If it’s

the same=ground state; if it’s different=excited state

• Warning: Both the formation of ions and the excited vs. ground state

involve electrons “doing things” but there are important differences

between the two.

Ions Excited State

• Definition: When an atom gains

or loses electrons and becomes

charged

• The amount of total electrons

changes

Example:

Na (atom) 2-8-1 Total e = 11

Na+1 (ion) 2-8 Total e = 10

Definition: When electrons absorb

energy and are found in a higher

energy configuration

• The amount of total electrons stay

the same, they just move shells!

(therefore still a neutral atom)

Example:

Na (ground state) 2-8-1 Total e = 11

Na (excited state) 2-7-2 Total e = 11

B. Bright Line Spectra

• When _________________________________, they __________________________ or an excited state.

o This is a very unstable/temporary condition

• ________________ rapidly ____ back down or drop ____________________(because they are unstable in the excited state)

• When excited electrons fall from an excited state to lower energy level (to the ground state), they _______________in the _____________(photons).

• One way this light is commonly analyzed is through a bright-line spectrum

o Recall, bright-line spectrum = a chart of lines of light given off when an electric current is run through an atom

ground state electrons absorb energy

jump to a higher energy level

Excited electrons fall

to a lower energy level

release energy form of light

• ________________________________________________

________________

o Fireworks are an example of this

• Spectra are unique for each element (like fingerprints are

unique for each person) so we can use

________________________________________

http://www.mhhe.com/physsci/chemistry/essentialchemis

try/flash/linesp16.swf

Different elements produce different colors of light or

different spectra

spectral lines to identify different elements

What elements are present in the mixture based on the

bright-line spectra?

Example:

Strontium and lithium

V. Isotopes• Isotopes = Atoms of the ______________but

______________________

o ________________ but __________________

Example:

1

1

1

0 1 11H

1 2 21H

2 3 31H

same # protons different # neutrons

same element different mass number

Practice

1. Determine the amount of each subatomic particle for

the following isotopes of Carbon (C-12, C-13, & C-14)

p = p = p =

n = n = n =

e = e = e =

6

6

6

6

7

6

6

8

6

*Notice-___________ have a ____________________

whereas _______ have a ______________________isotopes different # of neutrons ions different # of electrons

2. Two different isotopes of the same element must contain the

same number of

a. protons b. neutrons c. electrons

3. Two different isotopes of the same element must contain a

different number of

a. protons b. neutrons c. electrons

4. Isotopes of a given element have

a. the same mass number and a different atomic number

b. the same atomic number and a different mass number

c. the same atomic number and the same mass number

• We have learned that mass number is defined as the # of protons + the # of

neutrons.

• We have seen that mass numbers are all whole numbers. So what’s up with

the atomic mass given in the periodic table?

• Atomic mass and mass number are not the same; though they are similar

(the mass number should always be somewhat close to the atomic mass).

• So what is the atomic mass?

Atomic Mass = the ________________of

___________________________of an

element.

A weighted average takes in account

relative abundance, or

percentages/amount of each isotope

VI. Atomic MassA. Atomic Mass vs. Mass Number

weighted average all naturally occurring isotopes

• Yes, the atomic mass for each element is in the upper left-hand corner

on the periodic table. But how is it calculated?

Atomic Mass = the weighted average of an element’s naturally occurring

isotopes

(% abundance of isotope in decimal form) x (mass of isotope 1)

(% abundance of isotope in decimal form) x (mass of isotope 2)

+ (% abundance of isotope in decimal form) x (mass of isotope 3)

B. Calculating Atomic Mass

Mass Number Atomic Mass

The _____ of _________

of a given element.

The _____________of

___________of a given

element

mass one isotope average mass

all isotopes

Average Atomic Mass of the Element

Examples:

1. Carbon has two naturally occurring stable isotopes . 98.89% of

carbon atoms are C-12, while the remaining 1.108% are C-13.

What is the atomic mass of carbon?

Step 1: Convert % to decimal (by dividing by 100)

(0.9889)

Step 3: Add up the masses of isotopes

11.87 amu + 0.1440 amu =

(12 amu) = 11.87 amu

(0.01108) (13 amu) = 0.1440 amu

12.01 amu

Step 2: Multiply the decimal by its mass number

98.89%/100 = (0.9889) 1.108%/100 = (0.01108)

2. 92.21% of Si is found to be 27.98 amu, 4.70% is found to be

28.98 amu, and the remaining 3.09% is found to be 29.97.

Calculated the atomic mass of silicon

Step 1: Convert % to decimals (by dividing by 100)

Step 2: Multiply the decimals by the mass number

Step 3: Add up the masses of isotopes